[PPT]Organic Chemistry - Home | Rutgers University - Newarkandromeda.rutgers.edu/~frjordan/chapter1.ppt · Web viewWilliam H. Brown Christopher S. Foote Brent L. Iverson Covalent

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Organic Organic ChemistryChemistry

William H. BrownWilliam H. BrownChristopher S. FooteChristopher S. FooteBrent L. IversonBrent L. Iverson

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Covalent Covalent Bonding & Bonding & Shapes of Shapes of MoleculesMolecules

Chapter 1Chapter 1

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Organic ChemistryOrganic Chemistry The study of the compounds of carbon Over 10 million compounds have been identified• about 1000 new ones are identified each day!

C is a small atom • it forms single, double, and triple bonds• it is intermediate in electronegativity (2.5)• it forms strong bonds with C, H, O, N, and some metals

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Schematic View of an AtomSchematic View of an Atom• a small dense nucleus,

diameter 10-14 - 10-15 m, which contains positively charged protons and most of the mass of the atom

• an extranuclear space, diameter 10-10 m, which contains negatively charged electrons

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Electron Configuration of AtomsElectron Configuration of Atoms Electrons are confined to regions of space called

principle energy levels (shells)• each shell can hold 2n2 electrons (n = 1,2,3,4......)

Shell

Number of Electrons Shell

Can Hold

Relative Energiesof Electrons

in These Shells3218 8 2

4321

higher

lower

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Electron Configuration of AtomsElectron Configuration of Atoms Shells are divided into subshells called orbitals,

which are designated by the letters s, p, d, f,........• s (one per shell) • p (set of three per shell 2 and higher) • d (set of five per shell 3 and higher) .....

Shell Orbitals Contained in That Shell321 1s

2s, 2px, 2py, 2pz

3s, 3px, 3py, 3pz, plus five 3d orbitals

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Electron Configuration of AtomsElectron Configuration of Atoms Aufbau Principle:Aufbau Principle: • orbitals fill in order of increasing energy from lowest

energy to highest energy Pauli Exclusion Principle:Pauli Exclusion Principle: • only two electrons can occupy an orbital and their

spins must be paired Hund’s Rule:Hund’s Rule: • when orbitals of equal energy are available but there

are not enough electrons to fill all of them, one electron is added to each orbital before a second electron is added to any one of them

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Electron Configuration of AtomsElectron Configuration of Atoms The pairing of electron spins

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Electron Configuration of AtomsElectron Configuration of Atoms Table 1.3 The Ground-State Electron

Configuration of Elements 1-18

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Lewis Dot StructuresLewis Dot Structures Gilbert N. Lewis Valence shell:Valence shell: • the outermost occupied electron shell of an atom

Valence electrons:Valence electrons: • electrons in the valence shell of an atom; these

electrons are used to form chemical bonds and in chemical reactions

Lewis dot structure:Lewis dot structure: • the symbol of an element represents the nucleus and

all inner shell electrons• dots represent valence electrons

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Lewis Dot StructuresLewis Dot Structures Table 1.4 Lewis Dot Structures for Elements 1-18

N OB

H

Li Be

Na

He

Cl

F

S

Ne

Ar

C

SiAl P

1A 2A 3A 4A 5A 6A 7A 8A

Mg :

:::::

.

.

.

.

.

.

.

..

. .

. .

..

.

:

:

:

::::::::

::::::.

:::

:

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Lewis Model of BondingLewis Model of Bonding Atoms bond together so that each atom acquires

an electron configuration the same as that of the noble gas nearest it in atomic number• an atom that gains electrons becomes an anionanion• an atom that loses electrons becomes a cationcation• the attraction of anions and cations leads to the

formation of ionic solidsionic solids• an atom may share electrons with one or more atoms

to complete its valence shell; a chemical bond formed by sharing electrons is called a covalent bondcovalent bond

• bonds may be partially ionic or partially covalent; these bonds are called polar covalent bondspolar covalent bonds

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ElectronegativityElectronegativity Electronegativity:Electronegativity: • a measure of an atom’s attraction for the electrons it

shares with another atom in a chemical bond Pauling scalePauling scale• generally increases left to right in a row• generally increases bottom to top in a column

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Formation of IonsFormation of Ions A rough guideline: • ions will form if the difference in electronegativity

between interacting atoms is 1.9 or greater• example: sodium (EN 0.9) and fluorine (EN 4.0)• we use a single-headed (barbed) curved arrow to show

the transfer of one electron from Na to F

• in forming Na+F-, the single 3s electron from Na is transferred to the partially filled valence shell of F

Na + F Na+ F -••••••

••

••••

••

+ F-(1s22s22p6)Na+(1s22s22p6)F(1s22s22p5)+Na(1s22s22p63s1)

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Covalent BondsCovalent Bonds The simplest covalent bond is that in H2

• the single electrons from each atom combine to form an electron pair

• the shared pair functions in two ways simultaneously; it is shared by the two atoms and fills the valence shell of each atom

The number of shared pairs• one shared pair forms a single bond• two shared pairs form a double bond• three shared pairs form a triple bond

H H H-H+ • ΔH0 = -435 kJ (-104 kcal)/m ol•

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Polar and Nonpolar Covalent BondsPolar and Nonpolar Covalent Bonds Although all covalent bonds involve sharing of

electrons, they differ widely in the degree of sharing

We divide covalent bonds into• nonpolar covalent bonds• polar covalent bonds

Difference in ElectronegativityBetween Bonded Atoms Type of BondLess than 0.50.5 to 1.9Greater than 1.9

Nonpolar covalentPolar covalentIons form

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Polar and Nonpolar Covalent BondsPolar and Nonpolar Covalent Bonds• an example of a polar covalent bond is that of H-Cl• the difference in electronegativity between Cl and H is

3.0 - 2.1 = 0.9• we show polarity by using the symbols ++ and --, or by

using an arrow with the arrowhead pointing toward the negative end and a plus sign on the tail of the arrow at the positive end

H Cl+ δ-

H Cl

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Polar Covalent BondsPolar Covalent Bonds Bond dipole moment (Bond dipole moment (mm):): • a measure of the polarity of a covalent bond• the product of the charge on either atom of a polar

bond times the distance between the nuclei• Table 1.7 shows average bond dipole moments of

selected covalent bonds

H-OH-NH-C

H-S C-I

C-FC-ClC-Br C-N

-

C-O

C=N

C=O

-

Bond Dipole (D) Bond

1.4

1.51.51.41.2

Bond

Bond Dipole (D)

1.30.3 0.7

0.20.7

2.3

3.5

Bond Dipole (D)Bond

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Lewis StructuresLewis Structures To write a Lewis structure• determine the number of valence electrons• determine the arrangement of atoms• connect the atoms by single bonds• arrange the remaining electrons so that each atom has

a complete valence shell• show a bonding pair of electrons as a single line• show a nonbonding pair of electrons as a pair of dots• in a single bond atoms share one pair of electrons, in a

double bond they share two pairs of electrons, and in a triple bond they share three pairs of electrons

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Lewis Structures - Table 1.3Lewis Structures - Table 1.3

In neutral molecules• hydrogen has one bond• carbon has 4 bonds and no lone pairs• nitrogen has 3 bonds and 1 lone pair• oxygen has 2 bonds and 2 lone pairs• halogens have 1 bond and 3 lone pairs

H2O (8) NH3 (8) CH4 (8) HCl (8)

C2H4 (12) C2H2 (10) CH2O (12) H2CO3 (24)

H-O-H H-N-HH

H-C-HH

HH-Cl

H-C C-HH

HC O

H

HC C

H

H O OCH HO

Ethylene

Hydrogen chlorideMethaneAmmoniaWater

Carbonic acidFormaldehydeAcetylene

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Formal ChargeFormal Charge Formal charge:Formal charge: the charge on an atom in a

molecule or a polyatomic ion To derive formal charge

1. write a correct Lewis structure for the molecule or ion2. assign each atom all its unshared (nonbonding)

electrons and one-half its shared (bonding) electrons3. compare this number with the number of valence

electrons in the neutral, unbonded atom

Number of valence electrons in the neutral, unbonded atom

All unsharedelectrons

One half of all shared electrons

+Formalcharge

=

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Formal ChargeFormal Charge Example:Example: Draw Lewis structures, and show which atom in

each bears the formal chargeNH2

- HCO3- CO3

2-

CH3NH3+ HCOO- CH3COO-

(b) (c)

(d) (e) (f)

(a)

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Exceptions to the Octet RuleExceptions to the Octet Rule Molecules containing atoms of Group 3A

elements, particularly boron and aluminum

Aluminum chloride

:::

F B

F

F

Cl Al

Cl

Cl

6 electrons in the valence shells of boron

and aluminum

Boron trifluoride: :

: :

: :

::

::

:

::

::

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Exceptions to the Octet RuleExceptions to the Octet Rule Atoms of third-period elements have 3d orbitals

and may expand their valence shells to contain more than 8 electrons• phosphorus may have up to 10

::

Phosphorus pentachloride

Phosphoricacid

P

ClCl Cl

Cl ClCH3-P-CH3

CH3Trimethyl-phosphine

H-O-P-O-HO

O-H

:

:

:

::

::

::

::::

::

:

:

::

:

:

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Exceptions to the Octet RuleExceptions to the Octet Rule• sulfur, another third-period element, forms

compounds in which its valence shell contains 8, 10, or 12 electrons

:::H-S-H CH3-S-CH3 H-O-S-O-H

O

OO

Sulfuricacid

Hydrogensulfide

Dimethylsulfoxide

::

::

: : : :

: :

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Functional GroupsFunctional Groups Functional group:Functional group: an atom or group of atoms

within a molecule that shows a characteristic set of physical and chemical properties

Functional groups are important for three reason; they are1. the units by which we divide organic compounds into

classes2. the sites of characteristic chemical reactions3. the basis for naming organic compounds

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AlcoholsAlcohols contain an -OH (hydroxylhydroxyl) group

Ethanol may also be written as a condensed condensed structural formulastructural formula

H-C-C-O-HH

H

H

H::-C-O-H

Ethanol(an alcohol)

Functionalgroup

CH3-CH2-OH CH3CH2OHor

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AlcoholsAlcohols• alcohols are classified as primary (1°), secondary (2°),

or tertiary (3°) depending on the number of carbon atoms bonded to the carbon bearing the -OH group

CH3-C-OHCH3

CH3CH3-C-OH

CH3

H

CH3

CH3CH3-C-OH

A 1° alcohol A 2° alcohol A 3° alcohol

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AlcoholsAlcohols• there are two alcohols with molecular formula C3H8O

CH3CH2CH2OH

CH3CHCH3

OH

H-C-C-C-O-HH

H

H

H

H

H

C-C-C-HH

HOH

HHH

H

or

ora 2° alcohol

a 1° alcohol

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AminesAmines contain an amino groupamino group; an sp3-hybridized

nitrogen bonded to one, two, or three carbon atoms• an amine may by 1°, 2°, or 3°

CH3 N HH

CH3 N HCH3

CH3 N CH3CH3

Methylamine(a 1° amine)

Dimethylamine(a 2° amine)

Trimethylamine(a 3° amine)

: : :

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Aldehydes and KetonesAldehydes and Ketones contain a carbonyl (C=O) groupcarbonyl (C=O) group

C HO

CH3-C-HO

CH3-C-CH3

OCO

Functionalgroup

Acetaldehyde(an aldehyde)

Acetone(a ketone)

Functionalgroup

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Carboxylic AcidsCarboxylic Acids contain a carboxyl (-COOH) groupcarboxyl (-COOH) group

C OO

H CH3-C-O-HO

CH3COOH CH3CO2H: ::

: or orAcetic acid

(a carboxylic acid)Functional

group

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Carboxylic EstersCarboxylic Esters Ester:Ester: a derivative of a carboxylic acid in which

the carboxyl hydrogen is replaced by a carbon group

C OO

Functionalgroup

CH3-C-O-CH2-CH3

Ethyl acetate(an ester)

::

: :O

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Carboxylic AmideCarboxylic Amide Carboxylic amidearboxylic amide, commonly referred to as an

amideamide: a derivative of a carboxylic acid in which the -OH of the -COOH group is replaced by an amine

• the six atoms of the amide functional group lie in a plane with bond angles of approximately 120°

CH3-C-N-HH

Acetamide(a 1° amide)

OC NO

Functionalgroup

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VSEPRVSEPR Based on the twin concepts that• atoms are surrounded by regions of electron density• regions of electron density repel each other

HC C

H

O C

CH

NH

HC

HHO

H

CH C H

H

O

4 regions of e- density(tetrahedral, 109.5°)

3 regions of e- density(trigonal planar, 120°)

2 regions of e- density(linear, 180°)

H

CH H

HN

H HH

:

::

::

::

:

H HO

CH N

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VSEPR ModelVSEPR Model Example:Example: predict all bond angles for these molecules and

ions(a) NH4

+ (b) CH3NH2

(f) H2CO3 (g) HCO3-(e) CH3CH=CH2

(i) CH3COOH(h) CH3CHO

(d) CH3OH

(j) BF4-

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Polar and Nonpolar MoleculesPolar and Nonpolar Molecules To determine if a molecule is polar, we need to

determine • if the molecule has polar bonds• the arrangement of these bonds in space

Molecular dipole moment (Molecular dipole moment (mm):): the vector sum of the individual bond dipole moments in a molecule• reported in debyes (D)

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Polar and Nonpolar MoleculesPolar and Nonpolar Molecules these molecules have polar bonds, but each has

a zero dipole moment

O C O

Carbon dioxidem = 0 Δ

B

F

F

F

Boron trifluoriem = 0 Δ

C

Cl

ClClCl

Carbon tetrachloriem = 0 Δ

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Polar and Nonpolar MoleculesPolar and Nonpolar Molecules these molecules have polar bonds and are polar

moleculesN

HH

HO

H H

Waterm = 1.85Δ

Am m onia m = 1.47Δ

irectionof ipolem om ent

irectionof ipolem om ent

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Polar and Nonpolar MoleculesPolar and Nonpolar Molecules• formaldehyde has polar bonds and is a polar molecule

Formaldehyde m = 2.33 Δ

irectionof ipolem om ent H HC

O

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ResonanceResonance For many molecules and ions, no single Lewis

structure provides a truly accurate representation

Ethanoate ion(acetate ion)

CO

OH3CC

O

OH3C

-

-

and

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ResonanceResonance Linus Pauling - 1930s• many molecules and ions are best described by

writing two or more Lewis structures• individual Lewis structures are called contributing

structures• connect individual contributing structures by double-

headed (resonance) arrows• the molecule or ion is a hybrid of the various

contributing structures

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ResonanceResonance Examples:Examples: equivalent contributing structuresequivalent contributing structures

Acetate ion(equivalent contributing

structures)

Nitrite ion(equivalent contributing

structures)

:

:

:

:: : :

NO -

O: N

O

O -:

: : :

:

:

: :

: :C

O -

OCH3 C

O

O -CH3

: :

::

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ResonanceResonance Curved arrow:Curved arrow: a symbol used to show the

redistribution of valence electrons In using curved arrows, there are only two

allowed types of electron redistribution:• from a bond to an adjacent atom• from an atom to an adjacent bond

Electron pushing is a survival skill in organic chemistry• learn it well!

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ResonanceResonance All contributing structures must1. have the same number of valence electrons2. obey the rules of covalent bonding• no more than 2 electrons in the valence shell of H • no more than 8 electrons in the valence shell of a 2nd

period element• 3rd period elements, such as P and S, may have up to

12 electrons in their valence shells3. differ only in distribution of valence electrons; the

position of all nuclei must be the same4. have the same number of paired and unpaired electrons

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ResonanceResonance The carbonate ion, for example• a hybrid of three equivalent contributing structures• the negative charge is distributed equally among the

three oxygens

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ResonanceResonance Preference 1:Preference 1: filled valence shells• structures in which all atoms have filled valence shells

contribute more than those with one or more unfilled valence shells

••••••

Greater contribution; both carbon and oxygen have complete valence shells

Lesser contribution;carbon has only 6 electrons in its valence shell

+ +CCH3 OCH3 O

H

HC

H

H

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ResonanceResonance Preference 2:Preference 2: maximum number of covalent

bonds• structures with a greater number of covalent bonds

contribute more than those with fewer covalent bonds

••

••••

Greater contribution(8 covalent bonds)

Lesser contribution(7 covalent bonds)

+ +CCH3 OCH3 O

H

HC

H

H

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ResonanceResonance Preference 3:Preference 3: least separation of unlike charge• structures with separation of unlike charges contribute

less than those with no charge separation

Lesser contribution(separation of unlike

charges)

CH3-C-CH3 CH3-C-CH3

Greater contribution(no separation of unlike charges)

O -O

::

:::

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ResonanceResonance Preference 4:Preference 4: negative charge on the more

electronegative atom • structures that carry a negative charge on the more

electronegative atom contribute more than those with the negative charge on the less electronegative atom

CH3CH3H3CCH3H3CCO

CO

H3CC

O

(b)Greater

contribution

(c)Should notbe drawn

(a)Lesser

contribution

(1) (2)

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Quantum or Wave MechanicsQuantum or Wave Mechanics Albert Einstein: E = hn (energy is quantized)• light has particle properties

Louis deBroglie: wave/particle duality

Erwin Schrödinger: wave equation• wave function, wave function, : a solution to a set of equations that

depicts the energy of an electron in an atom• each wave function is associated with a unique set of

quantum numbers• each wave function occupies three-dimensional space

and is called an orbitalorbital• 22 is the probability of finding an electron at a given

point in space

l = hmν

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Shapes of 1Shapes of 1ss and 2 and 2ss Orbitals Orbitals Probability distribution (2) for 1s and 2s orbitals

showing an arbitrary boundary surface containing about 95% of the electron density

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Shapes of a Set of 2Shapes of a Set of 2pp Atomic Orbitals Atomic Orbitals Three-dimensional shapes of 2p atomic orbitals

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Molecular Orbital TheoryMolecular Orbital Theory Electrons in atoms exist in atomic orbitals Electrons in molecules exist in molecular orbitals

(MOs) Using the Schrödinger equation, we can

calculate the shapes and energies of MOs

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Molecular Orbital TheoryMolecular Orbital Theory Rules:• combination of n atomic orbitals (mathematically

adding and subtracting wave functions) gives n MOs (new wave functions)

• MOs are arranged in order of increasing energy• MO filling is governed by the same rules as for atomic

orbitals:• Aufbau principle: fill beginning with LUMO• Pauli exclusion principle: no more than 2e- in a MO• Hund’s rule: when two or more MOs of equivalent

energy are available, add 1e- to each before filling any one of them with 2e-

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Molecular Orbital TheoryMolecular Orbital Theory Terminology• ground state = lowest energy state• excited state = NOT lowest energy state• = sigma bonding MO• * = sigma antibonding MO• p = pi bonding MO• p* = pi antibonding MO• HOMO = highest occupied MO• LUMO = lowest unoccupied MO

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Molecular Orbital TheoryMolecular Orbital Theory Sigma 1s bonding and antibonding MOs

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Molecular Orbital TheoryMolecular Orbital Theory• MO energy diagram for H2: (a) ground state and (b)

lowest excited state

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Molecular OrbitalsMolecular Orbitals• computed sigma bonding and antibonding MOs for H2

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Molecular OrbitalsMolecular Orbitals pi bonding and antibonding MOs

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Molecular OrbitalsMolecular Orbitals• computed pi bonding and antibonding MOs for

ethylene

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Molecular OrbitalsMolecular Orbitals• computed pi bonding and antibonding orbitals for

formaldehyde

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Hybrid OrbitalsHybrid Orbitals The Problem:• bonding by 2s and 2p atomic orbitals would give bond

angles of approximately 90°• instead we observe bond angles of approximately

109.5°, 120°, and 180° A Solution• hybridization of atomic orbitals• 2nd row elements use sp3, sp2, and sp hybrid orbitals

for bonding

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Hybrid OrbitalsHybrid Orbitals Hybridization of orbitals (L. Pauling)• the combination of two or more atomic orbitals forms

a new set of atomic orbitals, called hybrid orbitals We deal with three types of hybrid orbitals

spsp33 (one s orbital + three p orbitals)spsp22 (one s orbital + two p orbitals)spsp (one s orbital + one p orbital)

Overlap of hybrid orbitals can form two types of bonds depending on the geometry of overlap bondsbonds are formed by “direct” overlappp bondsbonds are formed by “parallel” overlap

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spsp33 Hybrid Orbitals Hybrid Orbitals• each sp3 hybrid orbital

has two lobes of unequal size

• the sign of the wave function is positive in one lobe, negative in the other, and zero at the nucleus

• the four sp3 hybrid orbitals are directed toward the corners of a regular tetrahedron at angles of 109.5°

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spsp33 Hybrid Orbitals Hybrid Orbitals• orbital overlap pictures of methane, ammonia, and

water

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spsp22 Hybrid Orbitals Hybrid Orbitals• the axes of the three sp2 hybrid orbitals lie in a plane

and are directed toward the corners of an equilateral triangle

• the unhybridized 2p orbital lies perpendicular to the plane of the three hybrid orbitals

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Bonding in EthyleneBonding in Ethylene

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Bonding in FormaldehydeBonding in Formaldehyde

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spsp Hybrid Orbitals Hybrid Orbitals• two lobes of unequal size at an angle of 180°• the unhybridized 2p orbitals are perpendicular to each

other and to the line created by the axes of the two sp hybrid orbitals

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Bonding in Acetylene, CBonding in Acetylene, C22HH22

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Hybrid OrbitalsHybrid Orbitals

H-C-C-HH

H

H

H

H-C C-H

C CH

HH

H

OrbitalHybrid-ization

Types of Bonds

to Carbon Example

sp3 4 sigma bonds

sp2 3 sigma bondsand 1 pi bond

sp 2 sigma bondsand 2 pi bonds

Ethane

Ethylene

Acetylene

Name

PredictedBond

Angles

109.5°

120°

180°

GroupsBonded

to Carbon

4

2

2

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Bond Lengths and Bond StrengthsBond Lengths and Bond Strengths

H-C-C-HH

H

H

H

HC C

H

H H

H-C C-H

C-C

C-C

C-C

C-H

C-H

C-H

Formula BondOrbitalOverlap

Bond Length(pm)

Bond Strength[kJ (kcal)/mol]

sp3-sp3

sp2-1s

sp-sp, two 2p-2psp-1s

sp3-1s

sp2-sp2, 2p-2p

153.2111.4

133.9110.0

121.2109.0

376 (90)422 (101)

727 (174)464(111)

966 (231)556 (133)

Ethane

Ethylene

Acetylene

Name

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Covalent Covalent Bonds & Bonds & Shapes of Shapes of MoleculesMolecules

End Chapter 1End Chapter 1